The Periodic Table: an arrangement of elements in order of their atomic numbers so that elements with similar properties fall in the same column or group.
Chapter 5 The Periodic Table
By: Alec Root & Dennis Wagner
Edited by: Grace Pratt, Mana Aliabadi, and Megan Matthews
History of the Periodic Table
Important People
Dalton- proposed an atomic theory that said all atoms of the same element are the same and described the atom as a simple sphere.
was later proved wrong on both accounts
Dobereiner- discovered triads (groups of three elements that have similar behaviors)
Triads work because each element has the same number of valence electrons
Newlands- arranged elements by increasing atomic mass, saw that properties reoccurred every eight elements
Named the reoccurrence of properties every eight elements "Law of Octaves" ←
Came up with octet rule (an element will be stable with eight valence electrons)
Mendeleev- "Father of the Periodic Table"- created the first version of the periodic table, organized it by increasing atomic mass
Left spaces for elements that had yet to be discovered
Mosley- created modern periodic table, organized it by increasing atomic number
Wrote Periodic Law "physical and chemical properties of the elements are periodic functions of their atomic numbers"
Canizzaro- developed method for accurately measuring the realative masses of atoms
Allowed chemists to agree on standard values for atomic mass, initiated a search for relationships between an element's properties & atomic mass
The Periodic Table{an organized chart of the elements in order of increasing number of protons, aligned so each column contains elements w/ similar properties}
The Noble Gases
The noble gases are the las column in the periodic table (Group 18)
They consisit of Helium, Neon, Argon, Krypton, Xenon, and Radon
Only elements that are completely stable, (they do not need to lose/gain any electrons to become stable
Sometimes regarded as the most significant addition to the periodic table
The Lanthanides
They are the 14 elements with atomic numbers from 58 (Ce-Cerium) to 71 (Lu- Lutetium)
Are located in period 6 between groups 3 and 4
Their chemical and physical properties are so similar that separating them was a very tedious task requiring many chemists
The Actinides
The actinides are the 14 elements with atomic numbers from 90 (Th- Thorium) to 103 (Lr- Lawrencium)
Are located in period 7 between groups 3 and 4
Usually set off below the main portion of the periodic table to save space
Radioactive
Metals
Ductile- can be drawn into wires
Malleable- can be hammared into sheets
Conductors- transfer heat and electricity
Generally reflect light and have luster
Solids at room temperature with the exception of Hg
Non-Metals
Insulaters- do not transfer heat or electricity
Brittle and are not easily shaped or formed, break when hammered
Semi-Metals/ Metalloids
Elements along the division line between metals and non-metals
show properties of both metals and non-metals
Periodicity
The tendency of recurring patterns with respect to atomic numbers on the periodic table
The cause of periodicity is the arrangement of the electrons around the nucleus
Ex: the differances of the atomic numbers for noble gases is: 8, 8, 18, 18, 32.
Energy Sub Levels:
ENERGY LEVELS -- n=1, n=2, n=3....
ENERGY SUB-LEVELS
n=1 : 1 sub-level : spherical "S"
n=2 : 2 sub-levels : parabolic "S,P"
n=3 : 3 sub-levels : diffuse : "S,P,D"
n=4 : 4 sub-levels : fundamental : "S,P,D,F"
n=5 : 5 sub-levels : "S,P,D,F,G"
Each sub-level has several variations or "Orbitals"
Each orbital can hold a maximum of 2 electrons(s=1 orbital; p=3 orbitals; d=5 orbitals; f=7 orbitals)
ORBITAL DIAGRAMS:
1. Find which orbitals are filled
2. Construct Diagram
3. Add Electrons from bottom up
*Every orbital is represented by a dot*Electrons are represented by arrows.*Orbital D is always offset by 1.*Orbital F is always offset by 2.
Noble Gas Configurations:
Starting @ the element, move BACKWARDS to the nearest noble gas (column 18)
Write that gas in brackets { }
Write the rest of configuration normally.
Elements on the periodic are arranged so that they are with elements that share similar chemical properties.
Therefore they are organized vertically in groups, and horizontally in rows or periods.
The length of a row or period is determined by the number of electrons that occupy the sublevels being filled in the period.
The first period consists of the 1s sublevel being filled which can hold a total of two electrons. This means that only the elements Hydrogen and Helium are in the 1s sublevel.
In the second period, the 2s and 2p sublevels are being filled. The 1s sublevel can hold 2 electrons, and the 2p sublevel can hold 6 electrons, thus satisfying an octet.
Both the second and third period hold eight elements because they cover the 2s 2p and the 3s 3p sublevels.
The fourth and fifth periods hold 18 elements because they cover the 4s, 3d, 4p, and the 5s, 4d, 5p sublevels.
The seventh period has the same configuration as the sixth period, but only holds 29 elements because three of the elements have not yet been discovered.
When calculating what period an elements lands on, all you must do is see what the highest energy level is that the element occupies.
THE S-BLOCK ELEMENTS Group 1: "Alkali Metals"
Highly Reactive
Found freely in nature
Group 2: "Alkaline-earth metals"
Group two metals are found to be harder, denser, and stronger than the alkali metals in Group 1.
Too reactive to be found freely in nature.
Hydrogen and Helium
Hydrogen and Helium are two rare cases on the periodic table where their properties do not match their position on the table.
Even though Hydrogen has an electron configuration of 1s1 in the first energy level of the first group, it does not share the same properties of the other elements in the first group. In fact, Hydrogen’s properties do not resemble those any group on the periodic table.
Helium has a group configuration of ns2. The odd thing is, it is located in Group 18.
This is because its highest occupied energy level is filled by two electrons, meaning Helium experiences special chemical stability. Thus it is paired with the other unreactive elements of Group 18.
The d-Block Elements: Groups 3-12
The d-Block is capable of making 5 different orbitals. If two electrons are capable of fitting in each orbital, the d-block can hold 10 electrons.
The d-Block is odd because unlike the s, and p blocks, 1 energy level must be subtracted because it has a higher energy that the 4s sublevel. This means that the first level of the d-Block is 3d even though the d-Block begins in the fourth row.
The properties of d-Block metals are that they are metals with typical metallic properties and are often referred to as transition metals.
Some are not even capable of forming compounds freely in nature because they are so unreactive.
The p-Block Elements: Groups 13-18
The p-block elements are also called the main-group elements.
To find the number of total electrons in the highest occupied energy level for the p-block elements, you must take the group number minus 10.
At the far right end, it contains nometals with the exception of Helium and Hydrogen.
HALOGENS (Group 17)
They are the most reactive of the non-metals.
The high reactivity is caused by being one valence electron short of becoming stable.
They are sufficiently reactive in nature, but only in compound form.
However, as free metals they are stable in the presence of air.
The f-Block Elements: Lanthanides and Actinides
They are position in such a way because they must fill the 4f sublevel.
The Lanthanides are shiny metals with similar reactivity to the alkaline earth metals. They are the 14 elements from Lanthanum, La, to Hafnium, Hf in the sixth energy level.
Actinides are all radioactive.
Lanthanides
Actinides
Electron Configuration & Periodic Properties
Atomic Radii
An atomic radius is defined as one-half the distance between the nuclei of identical atoms that are bonded together.
Period Trends
As you move down the periodic table, more energy levels are added and thus the atomic radii increases
As you move across the periodic table, there are more protons in the nucleus that attract the electrons, making the atom smaller
Group Trends
In general, the atomic radii of the main-group elements increase down a group.
Ionization Energy
With enough energy, an electron can be removed from an atom.
An ion is an atom or group of bonded atoms that has a positive or negative charge.
Ionization is any process that results in the formation of an ion.
Ionization Energy is defined as the energy required to remove one electron from a neutral atom of an element.
As an element gets closer to becoming isoelectric with a noble gas, the harder it becomes to lose an electron.
Removing Electrons from Positive Ions
The energies for removal of additional electrons from an atom are referred to as the second ionization energy (IE2), third ionization energy (IE3), and so on.
Each successive electron removed from an ion feels an increasingly stronger effective nuclear charge because fewer electrons remain within the atom to shield the attractive force of the nucleus.
Electron Affinity
Electron affinity is the energy released when an atom gains an electron.
Period Trends
Halogens are the group that most readily accepts electrons.
This is why they are referred to as Love Electrons, because they want to become isoelectric with the nearest noble gas.
Ionic Radii
A positive ion is known as a cation. (remember the drawing of the cat streched out in a "+")
A negative ion is known as an anion. (remeber the drawing of an ant shaped like "---")
As there is a gradual increase of atomic radii down a group, there is also a gradual increase of ionic radii.
Valence Electrons
Valence electrons are the electrons available to be lost, gained, or shared in the formation of chemical compounds.
They are often located in incompletely filled main-energy levels.
The easiest way to determine how many valence electrons an atom has is to see what group it is in. Group 1 has 1 valence electron, group 2 has 2, Group 3 has 3, and so on excluding the d-block.
Electronegativity
Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound.
Electronegativity follows the same trend as Electron Affinity. It tends to (with some exceptions) increases going up and to the right of the periodic table.
Noble gases do not form compounds, and thus they do not have electronegativities.
The Periodic Table: an arrangement of elements in order of their atomic numbers so that elements with similar properties fall in the same column or group.
Chapter 5 The Periodic Table
By: Alec Root & Dennis Wagner
Edited by: Grace Pratt, Mana Aliabadi, and Megan Matthews
History of the Periodic Table
Important People
The Periodic Table{an organized chart of the elements in order of increasing number of protons, aligned so each column contains elements w/ similar properties}
Energy Sub Levels:
ORBITAL DIAGRAMS:
1. Find which orbitals are filled
2. Construct Diagram
3. Add Electrons from bottom up
*Every orbital is represented by a dot*Electrons are represented by arrows.*Orbital D is always offset by 1.*Orbital F is always offset by 2.
Noble Gas Configurations:
Elements on the periodic are arranged so that they are with elements that share similar chemical properties.
Therefore they are organized vertically in groups, and horizontally in rows or periods.
THE S-BLOCK ELEMENTS
Group 1: "Alkali Metals"
Group 2: "Alkaline-earth metals"
Hydrogen and Helium
The d-Block Elements: Groups 3-12
The p-Block Elements: Groups 13-18
The f-Block Elements: Lanthanides and Actinides
Actinides
Electron Configuration & Periodic Properties
Atomic Radii
Ionization Energy
Removing Electrons from Positive Ions
Electron Affinity
Period Trends
Ionic Radii
Electronegativity
Sources